Nitrogen trifluoride (NF3) is a colorless, toxic, odorless, nonflammable inorganic gas, with increasing use as an etchant in microelectronics, in the plasma etching of silicon wafers, in the cleaning of the PECVD chambers in the high volume production of liquid crystal displays and silicon-based thin film solar cells. In these applications NF3 is initially broken down in situ, by plasma. The resulting fluorine atoms are the active cleaning agents that attack the polysilicon and silicon oxide. NF3 has been considered as an environmentally preferable substitute for sulfur hexafluoride or perfluorocarbons such as hexafluoroethane.
NF3 can be prepared from the reaction of elemental fluorine (F2) with (i) ammonia (NH3), (ii) nitrogen trichloride (NCl3), or (iii) nonamethyltrisilylamine [(Me3Si)3N]. However, the reaction with NCl3 is not feasible because it is an explosive reagent at temperatures as low −196° C. The production of NF3 involve oxidation of N(3−) to N(3+), and this transformation in commercial processes has required the use of F2 as fluorinating agent. However, fluorine does not react with N2.
The reaction of fluorine gas (bp=−188° C.) with liquid NH3 (bp=−33° C.; mp=−78° C.) is thermodynamically favored, and proceed readily, resulting in N—H→N—F+H-F exchange, or simply N—H→N—N+H-F. However, the yield of NF3 gas is typically below 40%, because it is the least favored reaction among three possible mechanisms shown in Table 1.
Table 1: Thermodynamic feasibility of competitive mechanisms of the reactions of fluorine gas with liquid ammonia at −40° C.
TABLE 1Thermodynamic feasibility of competitive mechanisms of the reactions of fluorine gas with liquid ammonia at −40° C.mol F2/Log K,ΔH, KCal/Reactionmol NH3−40° C.mol1.5308−370.6 N2 (g) + 6 NH4F 2.0308−354.0 N2F2 (g) + 6 NH4F 3.0201−216.6 NF3 (g) + 3 NH4F
In 1903, Otto Ruff prepared nitrogen trifluoride by the electrolysis of a molten mixture of ammonium fluoride and hydrogen fluoride. Today, commercially viable high yielding processes for production of NF3 have improved the concentration of fluorine to the nitrogen based reagents, including ammonium bifluoride-hydrogen fluoride complexes in Table 2, and ammonium hexafluoride aluminates. These strategies have produced high yields for conversion of the ammonium salt to NF3, but have had accompanying high cost of electricity; poor turnover of the F2 flow in the process; and high maintenance cost from replacement of nickel reactor materials. Another approach has involved an engineering convenience by contacting independent gas phase dilutions of ammonia and fluorine in sulfur hexafluoride in a long vertical column at room temperature.
Table 2: Thermodynamic feasibility of competitive mechanisms of the reactions of fluorine gas with ammonium bifluoride
TABLE 2Thermodynamic feasibility of competitive mechanisms of the reactionsof fluorine gas with ammonium bifluoridemol F2/ΔH,molLog K,KCal/ReactionNH4F−140° C.mol1.5308−370.6 N2 (g) + 10 HF (I) 2.0308−354.0 N2F2 (g) + 10 HF (I) 3.0201−216.6 NF3 (g) + 5 HF (g)
Process economics show that the most expensive contributor to cost of production of NF3 has been fluorine, because three moles of fluorine would be required for every nitrogen atom in the most efficient reaction. Yet the most efficient process still generates significant quantities of nitrogen (N2) and tetrafluorohydrazine (N2F2), and requires costly purification procedures. Thus, the cost of commercial production of NF3 is relatively higher than it is for many other inorganic fluorides.
Thermodynamic feasibility of the production of NF3 versus N2F2/N2 was highest at lower temperatures. The ratios of the logarithm of equilibrium constants for competitive mechanisms in the reaction of fluorine with NH4F is illustrated in Table 3.
Table 3: Comparison of temperature dependent Log k of equilibrium constant among three possible mechanisms shown in Table 1
TABLE 3Comparison of temperature dependent Log k of equilibriumconstant among three possible mechanisms shown in Table 1−80° C.−40° C.−0° C.20° C.140° C.N2/NF3N2F2/NF3N2/NF3N2F2/NF3N2/NF3N2F2/NF3N2/NF3N2F2/NF3N2/NF3N2F2/NF3308/219308/219308/184302/184283/159263/159266/149247/149202/110186/1101.411.411.671.641.781.651.791.661.841.69*Data obtained from the HSC Chemistry 7.0 software
The present invention relates to using trimethylsilylamines reagents to prepare NF3. The trimethylsilyl-nitrogen bond undergoes facile cleavage in the presence of reactive fluoride ion to produce very strong trimethylsilyl-fluoride bond, and the substitution of nitrogen by the conjugate anion. When the reagent is fluorine, an N—F bond is formed. Two trimethylsilylamines can be used to produce NF3 in commercial production, that is tris(trimethylsilyl)amine (also called nonatrimethyltrisilazane) and hexamethyldisilazane. The trimethylsilylamines are very soluble in solvents that can be used as a medium for low temperature liquid processes, such as acetonitrile, and fluorocarbons. With this, a careful study of the temperature-dependent kinetics of the process can be determined to accurately control the effective supply of fluorine to the process.